Sodium bicarbonate
Sodium bicarbonate, with the chemical formula NaHCO₃ and commonly known as baking soda or bicarbonate of soda, is the sodium salt of carbonic acid, appearing as an odorless white crystalline powder with a slightly alkaline taste.[1] It occurs naturally as the mineral nahcolite in certain mineral springs and deposits, and is produced industrially on a large scale by reacting carbon dioxide with solutions of sodium carbonate derived from the Solvay process or processed from trona ore.[1][2] The compound decomposes upon heating above 50°C to yield sodium carbonate, water, and carbon dioxide, a reaction central to its utility as a leavening agent in baking where it generates gas to raise dough.[1] It is highly soluble in water (approximately 96 g/L at 20°C) but insoluble in ethanol, enabling its roles in aqueous pH buffering and cleaning applications.[1] In household and industrial contexts, sodium bicarbonate functions as a mild abrasive cleaner, deodorizer, and component in fire extinguishers due to its ability to release CO₂ to smother flames.[1] Medically, it is employed as an antacid to neutralize excess gastric acid and intravenously to correct metabolic acidosis, hyperkalemia, and certain toxicities by restoring acid-base balance, though excessive use risks alkalosis or gastrointestinal issues.[1][3] Overall, its low toxicity (oral LD50 around 4,220 mg/kg in rats) and versatility make it a staple in food, pharmaceutical, and chemical sectors, with global production exceeding millions of tons annually.[1]Chemical Identity and Properties
Molecular Structure and Composition
Sodium bicarbonate, with the chemical formula NaHCO₃, is an ionic salt composed of one sodium cation (Na⁺) and one bicarbonate anion (HCO₃⁻).[4][5] The compound's molar mass is 84.0066 g/mol, reflecting the atomic masses of its constituent elements: sodium (22.9898 g/mol), hydrogen (1.00794 g/mol), carbon (12.0107 g/mol), and three oxygen atoms (each 15.9994 g/mol).[5][6] The bicarbonate anion (HCO₃⁻) consists of a central carbon atom bonded to three oxygen atoms in a trigonal planar arrangement, with one oxygen atom protonated to form a hydroxyl group (-OH).[7] This geometry arises from the sp² hybridization of the carbon atom, resulting in bond angles near 120° around the carbon center.[8] The two non-protonated carbon-oxygen bonds exhibit resonance, delocalizing the negative charge and imparting partial double-bond character, with typical bond lengths of approximately 1.25–1.30 Å for C-O and 1.36–1.43 Å for C=O-like bonds.[7] In the solid state, sodium bicarbonate forms a crystalline lattice where each sodium cation is coordinated to six oxygen atoms from surrounding bicarbonate anions, while each bicarbonate anion links to multiple sodium cations via its oxygen atoms.[4] This ionic assembly accounts for the compound's stability and solubility properties, distinct from covalent molecular structures.[4]Physical Characteristics
Sodium bicarbonate manifests as an odorless white crystalline powder, granular solid, or lumps, forming colorless or white monoclinic prisms.[4] It exhibits a slightly alkaline, salty taste.[9] The true density of the crystalline solid measures 2.20 g/cm³, while bulk powder densities vary from 1.1 to 1.3 g/cm³ based on particle characteristics.[10][4] The compound adopts a monoclinic crystal system, with sodium cations coordinated to bicarbonate anions in a structured lattice.[4] Solubility in water is temperature-dependent, registering 69 g/L at 0 °C, 96 g/L at 20 °C, and 165 g/L at 60 °C; it is practically insoluble in alcohol.[11] Sodium bicarbonate remains stable in dry air but undergoes gradual decomposition in moist environments due to interaction with atmospheric water vapor and carbon dioxide.[4] Thermal behavior involves no distinct melting point; decomposition initiates above 50 °C, evolving carbon dioxide, water vapor, and sodium carbonate, with full conversion to sodium carbonate occurring near 270 °C.[4] This endothermic process underpins its utility in applications requiring gas release upon heating.[4]
Chemical Reactivity and Stability
Sodium bicarbonate exhibits high chemical stability under standard ambient conditions, remaining non-reactive when stored in cool, dry environments within airtight containers.[12] It maintains integrity for at least two years at 25 °C and 60% relative humidity, though prolonged exposure to elevated humidity above 76% at 25 °C or 48% at 40 °C can induce gradual moisture sorption and potential decomposition.[13][14] Upon heating, sodium bicarbonate undergoes thermal decomposition via the endothermic reaction $2 \mathrm{NaHCO_3} \rightarrow \mathrm{Na_2CO_3} + \mathrm{H_2O} + \mathrm{CO_2}, initiating gradually above 80 °C (176 °F) and accelerating with temperature increase up to 180 °C, where the process follows pseudo-first-order kinetics influenced by partial pressure of CO₂.[15][16] This decomposition yields sodium carbonate, water vapor, and carbon dioxide gas, with no explosive hazard under controlled conditions due to the solid-state nature of pure samples. In aqueous solutions, sodium bicarbonate dissociates partially to form a mildly alkaline buffer (pH ≈ 8.3), acting as the conjugate base of carbonic acid; stability in polyolefin bags persists for 48 hours at 50 meq/L room temperature but shortens for higher concentrations (30 hours at 150 meq/L).[4] Reactivity with acids proceeds via protonation of the bicarbonate ion: \mathrm{NaHCO_3 + H^+ \rightarrow Na^+ + H_2O + CO_2}, rapidly liberating carbon dioxide gas, as observed in gastric neutralization where excess bicarbonate beyond acid equivalence remains unreacted.[4][17] Toward strong bases, sodium bicarbonate behaves as a weak acid, undergoing the reaction \mathrm{NaHCO_3 + NaOH \rightarrow Na_2CO_3 + H_2O}, which shifts equilibrium to carbonate formation without gas evolution.[18] It shows no significant reactivity with oxygen or neutral conditions, underscoring its utility in fire suppression via CO₂ release upon moist acid contact, though pure dry forms do not ignite or propagate flames.[19][17]Production and Sourcing
Industrial Synthesis via Solvay Process
The Solvay process, patented by Ernest Solvay in 1861 and first commercialized in 1865, synthesizes sodium bicarbonate (NaHCO3) as a key intermediate from sodium chloride (brine), ammonia (NH3), carbon dioxide (CO2), and water.[20][21] The process exploits the low solubility of NaHCO3 in concentrated brine (approximately 9 g/100 mL at 20°C), allowing its selective precipitation while sodium chloride remains in solution./05:_Chemical_Reactions-_Making_Materials_Safely_and_Sustainable/5.13:Industrial_Chemical_Reactions-_The_Solvay_Process) The process begins with the calcination of limestone (CaCO3) in kilns at temperatures around 900–1000°C to produce quicklime (CaO) and CO2:CaCO3 → CaO + CO2.[22] The CO2 is captured and purified for subsequent use. Separately, saturated brine (typically 20–25% NaCl aqueous solution) is saturated with gaseous ammonia under controlled conditions to form ammoniated brine.[23] In the carbonation tower, CO2 is bubbled through the ammoniated brine at pressures of 2–3 atm and temperatures maintained below 40°C to precipitate NaHCO3 via the reaction:
NaCl + NH3 + CO2 + H2O → NaHCO3 ↓ + NH4Cl./05:_Chemical_Reactions-_Making_Materials_Safely_and_Sustainable/5.13:Industrial_Chemical_Reactions-_The_Solvay_Process)[22] This step occurs in multi-stage towers where CO2 absorption drives the equilibrium toward bicarbonate formation, with the precipitate forming a slurry that is filtered using rotary vacuum filters or centrifuges. The solid NaHCO3 is then washed with cold water to remove impurities like NH4Cl and dried at low temperatures (around 100–150°C) to yield a purity exceeding 99% for commercial grades.[24] Ammonia is recovered for recycling by reacting the mother liquor (containing NH4Cl) with CaO (from limestone calcination) in a slaker:
2NH4Cl + CaO → CaCl2 + H2O + 2NH3, followed by distillation to reclaim gaseous NH3, achieving near-complete ammonia recovery (over 98%) and minimizing raw material losses./05:_Chemical_Reactions-_Making_Materials_Safely_and_Sustainable/5.13:Industrial_Chemical_Reactions-_The_Solvay_Process) Calcium chloride byproduct is typically discharged as wastewater or used in de-icing, though modern plants incorporate waste minimization. While the process is energy-intensive (requiring about 1.5–2 tons of steam per ton of NaHCO3 equivalent), its efficiency stems from raw material abundance and byproduct recycling, making it economically viable for large-scale production.[25][23] In contemporary operations, NaHCO3 isolation bypasses the traditional calcination to sodium carbonate, enabling direct marketing of bicarbonate for applications like food processing and pharmaceuticals.[24]